How are metals extracted from their ores? How are metals prevented from reverting to their oxides? Reactivity series Electrolysis background How to extract metals IRON Preventing iron reverting BRINGS YOU BACK HERE ALUMINIUM Preventing aluminium reverting Copper purification END

Some metals are more reactive than others. A list of metals in order of their reactivity is called the REACTIVITY SERIES. Reactions with air / oxygen HIGH REACTIVITY LOW Potassium (K) Sodium (Na) Calcium (Ca) Magnesium (Mg) Aluminium (Al) Zinc (Zn) Iron (Fe) Tin (Sn) Lead (Pb) Copper (Cu) Silver (Ag) Gold (Au) Platinum (Pt) All these metals burn vigorously with a bright flame. copper + oxygen  copper oxide The metal burns with a bright, white flame, leaving a white ash – magnesium oxide. 2 Cu + O2  CuO  balanced 2 NOT balanced The burning of magnesium ….. 2 Mg + O2  MgO The metal slowly gets a black covering of copper oxide when heated.  balanced An unreactive element, found in the earth’s crust as the metal itself. Little reaction with oxygen.

IT ALL FITS IN WITH THE REACTIVITY SERIES SUMMARY OF REACTIONS WITH WATER The less reactive metals do nothing. Magnesium reacts only when heated with steam. Calcium reacts gently when cold Sodium & potassium react increasingly violently when cold. IT ALL FITS IN WITH THE REACTIVITY SERIES Reactions of metals with water HIGH REACTIVITY Potassium (K) Sodium (Na) Calcium (Ca) Magnesium (Mg) Aluminium (Al) Zinc (Zn) Iron (Fe) Lead (Pb) Copper (Cu) Gold (Au) Calcium slowly dissolves in cold water and fizzes. The solution is an ALKALI The gas is hydrogen. Sodium and potassium run around on the surface of cold water, produce heat & dissolve, with lots of fizzing. The solution is a strong ALKALI The gas is hydrogen. The hot metal burns with a bright, white flame when steam is passed over it. It produces magnesium oxide. Hydrogen can be detected.. No reactions. magnesium (s) + water (g)  + hydrogen(g) magnesium oxide (s) calcium(s) + water(l)  + hydrogen(g) calcium hydroxide(aq) sodium (s) + water (l)  sodium hydroxide(aq) + hydrogen(g) This state symbol, (g) for “gas” tells us it’s STEAM  balanced 2 Na + H2O  NaOH + H2 Mg + H2O  MgO + H2 2 2  balanced Note: metals that react with STEAM produce the OXIDE metals that react with WATER produce the HYDROXIDE liquid water alkalis contain the hydroxide ion, OH-  balanced Ca + H2O  Ca(OH)2 + H2 2 NOT balanced NOT balanced

Put some zinc metal into blue copper sulphate solution. A MORE REACTIVE METAL CAN DISPLACE A LESS REACTIVE ONE FROM ITS COMPOUNDS MORE REACTIVE Sodium (Na) Calcium (Ca) Magnesium (Mg) Aluminium (Al) Zinc (Zn) Iron (Fe) Lead (Pb) Copper (Cu) Gold (Au) Put some zinc metal into blue copper sulphate solution. The zinc displaces the copper and forces the copper to become an element. Because zinc is more reactive, it has a greater tendency to become a compound than copper has. NOTE : Lead, for example would not displace iron from iron chloride solution because lead is LESS reactive. Pb(s) + FeCl3 (aq) x LESS REACTIVE Zinc(s) + copper sulphate(aq)  zinc sulphate(aq) + copper(s) In the test tube we see the zinc dissolving, the copper sulphate solution getting less blue and copper metal appearing at the bottom. copper displaced from its compound zinc is more reactive & becomes a compound  balanced Zn + CuSO4  ZnSO4 + Cu REMEMBER The MORE reactive a metal is, the more it likes being in a COMPOUND The LESS reactive a metal is, the more it likes being an ELEMENT The zinc has displaced the copper from its compound.

x  2 2 Al + Fe2O3  Al2O3 + Fe Cu(s) + ZnO (s)  Sodium (Na) A MORE REACTIVE METAL CAN DISPLACE A LESS REACTIVE ONE FROM ITS OXIDE MORE REACTIVE Sodium (Na) Calcium (Ca) Magnesium (Mg) Aluminium (Al) Zinc (Zn) Iron (Fe) Lead (Pb) Copper (Cu) Gold (Au) Heat aluminium powder with powdered iron(III) oxide. The mixture flares up. The aluminium is displacing the iron from its oxide. The aluminium is taking the oxygen from the iron, leaving iron metal. LESS REACTIVE NOTE : There would be no reaction between copper and zinc oxide because copper is LESS reactive. Cu(s) + ZnO (s)  x Aluminium(s) + iron oxide(s)  + iron(s) Aluminium oxide(s)  balanced 2 Al + Fe2O3  Al2O3 + 2 Fe NOT balanced

REACTIONS OF METALS WITH ACIDS MORE REACTIVE REACTIONS OF METALS WITH ACIDS Sodium (Na) Calcium (Ca) Magnesium (Mg) Aluminium (Al) Zinc (Zn) Iron (Fe) Lead (Pb) Hydrogen (H) Copper (Cu) Gold (Au) React with cold water : violent with acids Less reactive than hydrogen. Cannot displace it. These metals have no reaction with acid. A more vigorous reaction. Hot solution & lots of bubbles. A reaction of moderate speed. Steady production of bubbles. Acids are compounds containing hydrogen as the H+ ion. We must include hydrogen in the reactivity series. LESS REACTIVE Metal + acid  salt + hydrogen Metals that are more reactive than hydrogen, displace it from the acid and set the hydrogen free as the gaseous element. Zinc + sulphuric acid  zinc sulphate + hydrogen (s) (aq) (g)

   metal(s) + acid(aq) salt(aq) +hydrogen(g) 2 2 Examples iron(II) sulphate iron + sulphuric acid  + hydrogen  balanced H2SO4  Fe + FeSO4 + H2 magnesium chloride hydrochloric acid magnesium +  + hydrogen  balanced Mg + 2 HCl  MgCl2 + H2  balanced lead nitrate NOT balanced lead + nitric acid  + hydrogen H2 Pb + 2 HNO3  Pb(NO3) 2 + NOT balanced

EXTRACTING METALS FROM THEIR ORES. The earth’s crust contains metals and metal compounds, always found mixed with other substances. Gold is an unreactive metal, at the bottom of the reactivity series. It is found in the earth as the metal itself and chemical separation is not needed. In ORES, the metal or its compound is concentrated enough to make it economic to extract the metal. Often an ore contains a metal oxide or something that can easily be changed into a metal oxide. The ore haematite is iron(III) oxide, Fe2O3 The ore bauxite is aluminium oxide, Al2O3 To extract the metal, the oxygen must be removed from the metal oxide. ThIs process is called REDUCTION.

HOW A METAL IS EXTRACTED FROM ITS ORE DEPENDS ON HOW REACTIVE IT IS. MORE REACTIVE Hydrogen will displace less reactive metals from oxides of those metals. Because its more reactive, it is able to remove the oxygen and combine with it to make water. Aluminium (Al) Zinc (Zn) Iron (Fe) Lead (Pb) Hydrogen (H) Copper (Cu) Silver (Ag) Gold (Au) So hydrogen can reduce ; copper oxide silver oxide Pass a stream of hydrogen gas over heated, black copper oxide. It will gradually go brown as copper metal is made. Hydrogen cannot reduce ; LESS REACTIVE lead oxide iron oxide zinc oxide, etc copper oxide(s) + hydrogen(g) = copper(s) + water(g) H2  CuO + Cu + H2O

Electricity generation is a modern development. HOW CARBON CAN BE USED TO EXTRACT A METAL FROM ITS ORE. (this is an important, central point) Using first wood and then coal, this has been known about for thousands of years. (since the Iron Age) The difficulty of extracting aluminium from its oxide meant a stronger method, electrolysis, was needed & this is expensive. Electricity generation is a modern development. So we have only been able to extract the more reactive metals in recent times. MORE REACTIVE Sodium (Na) Calcium (Ca) Magnesium (Mg) Aluminium (Al) Carbon (C) Zinc (Zn) Iron (Fe) Lead (Pb) Copper (Cu) Gold (Au) Aluminium is MORE reactive than carbon and so carbon does NOT have the ability to remove oxygen from aluminium oxide. This is the position of the non-metal, carbon in the reactivity series. or zinc A metal such as iron, which is less reactive than carbon CAN be extracted from its ore using carbon. Because it is more reactive, carbon can take the oxygen from the iron oxide. As a result, iron is obtained from its ore using carbon in the Blast Furnace. Aluminium has to be extracted from Bauxite using ELECTROLYSIS. These two methods are considered in the following slides. LESS REACTIVE

EXTRACTION OF IRON IN THE BLAST FURNACE The raw materials are : coke (C) haematite (Fe2O3) limestone (CaCO3) EXTRACTION OF IRON IN THE BLAST FURNACE 1. The coke burns in an exothermic reaction. Energy is released, raw material gets hot and carbon dioxide is formed. waste gases 2. At the high temperatures, the carbon dioxide reacts with more coke to make carbon monoxide. 3. The carbon monoxide REDUCES the iron oxide in the ore, into molten iron which flows to the bottom of the furnace. carbon + oxygen  carbon dioxide C + O2  CO2 carbon dioxide monoxide + carbon  4. Limestone is added to remove acidic impurities forming a molten slag which floats on top of the iron. Hot air is blown in CO2 + C  CO 2  balanced NOT balanced SLAG MOLTEN IRON

The process of removing oxygen from the ore is called REDUCTION. Hot air MOLTEN IRON SLAG The carbon monoxide REDUCES the iron oxide in the ore, into molten iron which flows to the bottom of the furnace. This is the central, important reaction in the Blast Furnace. The process of removing oxygen from the ore is called REDUCTION. The carbon monoxide combines with the oxygen from the iron ore to produce carbon dioxide. This is called OXIDATION. Iron(III) oxide +  iron + carbon monoxide dioxide This is OXIDATION This is REDUCTION Fe2O3 + CO  Fe + CO2 3  balanced 2 3 NOT balanced

HOW CAN IRON BE PREVENTED FROM REVERTING TO ITS OXIDE ? Iron & steel corrode with air and water more quickly than most transition metals. It can be prevented by : connecting iron with a more reactive metal. “Galvanised iron” is the name given to iron plated with zinc. The more reactive metal corrodes first, (zinc) making sure the iron does NOT corrode. That is why the zinc would be called a “sacrificial metal”. magnesium wire underground iron pipe Here, the more reactive magnesium is being sacrificed, saving the iron pipe. Steel can be mixed with other metals such as chromium (Cr) to make an alloy which will not rust. Lots of cutlery made in Sheffield is manufactured from this “stainless steel”. Theory of Electrolysis ALUMINIUM

- MAKING ALUMINIUM METAL INDUSTRIALLY USING ELECTROLYSIS. + This large block of carbon is connected to the positive of the voltage supply. The ions are attracted as shown Oxygen forms at the positive electrode MOLTEN ALUMINIUM OXIDE IN CRYOLITE O2- Al3+ - The lining of this large tank is made of carbon and is connected to the negative of the voltage supply. The ions must be able to move but melting the aluminium oxide is difficult because of its high melting point. So the aluminium oxide is dissolved in molten aluminium compound called CRYOLITE at a much lower temperature. The raw material from which aluminium is produced is aluminium oxide, Al2O3 which contains the ions Al3+ and O2- The ore is called BAUXITE which needs purifying to get the aluminium oxide. The aluminium forms at the negative electrode The oxide ion loses electrons and is oxidised : 2O2- - 4e  O2 This oxygen reacts with the carbon (+) electrode and makes carbon dioxide gas. So the electrode burns away quickly and often has to be replaced. The aluminium ion gains electrons and is reduced Al3+ + 3e  Al

HOW CAN ALUMINIUM BE PREVENTED FROM REVERTING TO ITS OXIDE ? Aluminium is almost as reactive as magnesium & we might expect it to corrode quickly over time, or to burn up (oxidise) completely if it got hot. But it does not do this. Gets covered in a thin layer of aluminium oxide as soon as it is made. Aluminium metal This surface layer forms a barrier to oxygen and water & so prevents further corrosion. Aluminium is a useful structural metal, low in density. It can be made harder, stronger & stiffer by mixing it will small amounts of other metals such as magnesium, to make an alloy.

Cu - e  Cu2+ Cu2+ + e  Cu 2 PURIFICATION OF COPPER - - - - - 2 - + + The negative electrode is pure copper. + is the copper ion, Cu2+ The solution must contain the Cu2+ion Copper sulphate(aq) will do. The positive electrode is scrap copper containing impurities. - + The scrap copper dissolves in the solution. Pure copper gets plated onto the negative electrode. - - + - + - + - + + - + - - impurities separated out + - + OXIDATION AT THE POSITIVE ELECTRODE 2 Cu - e  Cu2+ REDUCTION AT THE NEGATIVE ELECTRODE Cu2+ + e  Cu 2

THE END These are the remains of some of the earliest Blast Furnaces. (Coalbrookdale Museum of Iron) This is a painting of the “Bedlam” Blast Furnaces at night, hundreds of years ago. To find out more about the Steel Industry, visit : The Ironbridge Gorge Museums, Telford Kelham Island Museum, Sheffield

Electrolysis

What sort of substances conduct electricity? ALL METALS This is because they have some FREE ELECTRONS in them that can move throughout the metal These electrons can carry the current because: They have an electric charge They can move GRAPHITE (a form of carbon) Graphite is the ONLY non-metal element that conducts electricity Like metals, it conducts because it has FREE ELECTRONS IONIC SUBSTANCES when MOLTEN or in SOLUTION When ionic substances are molten (= in liquid form) or in solution, their positively and negatively charged ions can move – so they conduct. When they are solid, the ions can’t move – so no current flows,

When we pass a current through a metal or graphite, the substance is not changed. When we pass a current through an ionic substance that’s molten or in solution, CHEMICAL CHANGES happen. This is ELECTROLYSIS

What we need for electrolysis + - e- ELECTRODES to carry the current into the solution or molten compound These are made of graphite or a metal Electrical wires to make the circuit A cell or battery e- e- e- e- The ELECTROLYTE – that’s the solution or molten ionic compound the electrodes are dipped in. The cell pushes electrons around the circuit FROM its NEGATIVE terminal TO its POSITIVE terminal This means that the positive electrode ends up short of electrons the negative electrode ends up with extra electrons YES – that IS the right way round – even though it’s opposite to the way we think of the current going…

- - + + Opposite charges attract negative ion, Cl- + positive ion, Cu2+ Opposite charges attract The negative ions move to the positive electrode The positive ions to the negative electrode The ionic compound is decomposed by the passage of electric current. copper chloride(aq) Elements are released at the electrodes as gases or metals. In this example, you would get copper metal made at the negative electrode and chlorine gas appearing at the positive electrode.

Here are Cu2+ ions moving to the negative electrode. Positive ions in the solution are attracted to negative electrode – opposite charges attract. When electrons are gained by a positive ion, the name of the chemical change is REDUCTION. REDUCTION IS THE GAIN OF ELECTRONS. THE COPPER ION HAS BEEN REDUCED 2e- + Cu2+  Cu The electrode is negative because it has too many electrons As they get close, the ions gain electrons from the electrode and the Cu2+ is neutralised. e- go to ion Cu This makes copper the element, which covers the electrode. e- go to ion. TWO ELECTRONS FROM THE CATHODE A NEUTRAL ATOM OF THE ELEMENT COPPER. ARE ADDED TO THE COPPER ION

When electrons are lost by a negative ion, the name of the chemical change is OXIDATION. OXIDATION IS THE LOSS OF ELECTRONS THE CHLORIDE ION HAS BEEN OXIDISED e- go to cell Cl- This is what happens at the positive electrode when chloride ions, Cl- are present in the electrolyte Here are negative chloride ions attracted towards the positive electrode. Opposite charges attract. This electrode is positive because some electrons have been removed by the cell. As they get close, each Cl- ion loses an electron which goes onto the electrode. The ion becomes electrically neutral the ion loses an e- Cl2 We have made chlorine the element. The neutral atoms join in pairs to make chlorine molecules, Cl2 which bubble off as a gas. 2Cl- - 2e  Cl2 TWO CHLORIDE IONS, EACH WITH AN EXTRA ELECTRON THE TWO ELECTRONS LOST BY THE IONS GO TO THE ELECTRODE A NEUTRAL CHLORINE MOLECULE END